Grade 11: Unit 6- Energetics

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IB Chemistry HL1

• Grade 11 Unit 6- Energetics
• IB Topic 5

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System and Surroundings

• The system is the name we give the sample or
  reaction vessel of interest.
• The surroundings are everything else in the
  universe.
• When a chemical change happens in an open system
  matter and energy can be exchanged between the
  system and the surroundings.
• In a closed system only energy can be exchanged
  with the surroundings.

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System and Surroundings
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Energy

• Energy is defined as the ability to do work.
  Energy is often converted from one form to
  another during physical and chemical changes.
• Thermochemistry is the study of energy changes
  associated with chemical reactions.
• Chemical energy is the energy stored in chemical
  bonds. It is a type of potential energy.

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Thermochemistry

• Most reactions absorb or evolve energy usually in
  the form of heat but chemical reactions can also
  produce light, electricity and mechanical energy
  used to do work.
• Energy is measured in joules, J.
• 1000 J 1 kJ.
• Physical changes like change of state/phase also
  have heat energy changes.

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Exothermic and Endothermic Reactions

• 5.1.1 Define the terms exothermic reaction,
  endothermic reaction and standard enthalpy change
  of reaction (?H).
• Standard enthalpy change is the heat energy
  transferred under standard conditionspressure
  101.3 kPa, temperature 298 K. Only ?H can be
  measured, not H for the initial or final state of
  a system.

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Exothermic and Endothermic Reactions

• 5.1.3 Apply the relationship between temperature
  change, enthalpy change and the classification of
  a reaction as endothermic or exothermic.
• 5.1.4 Deduce, from an enthalpy level diagram, the
  relative stabilities of reactant and products,
  and the sign of the enthalpy change for the
  reaction.

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Enthalpy

• Enthalpy is the total energy of a system, some of
  which is stored as chemical potential energy in
  the chemical bonds.
• Enthalpy is given the symbol H.
• Enthalpy is also known as the heat content of a
  system.
• We cannot measure the enthalpy content of a
  system but we can measure changes in it.

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Enthalpy Change

• In chemical reactions, bonds are broken and made,
  but the energy absorbed breaking bonds is almost
  never exactly equal to that released in making
  new bonds.
• All reactions are accompanied by a change in the
  potential energy of the bonds and hence an
  ENTHALPY CHANGE.
• There is no absolute zero for enthalpy so
  absolute enthalpies cannot be measured only the
  change in enthalpy that occurs during a reaction.

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Enthalpy Change

• The enthalpy change of a reaction is given the
  symbol ?H.
• ?H is the difference in the enthalpy between the
  products and the reactants.
• ?H H(products) - H(reactants) when at constant
  pressure.
• Enthalpy level diagrams are used to show the
  change in enthalpy of a system during a change.

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Enthalpy Level Diagram
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Enthalpy Level Diagram
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Endothermic Reaction

• An endothermic reaction is one where energy is
  transferred from the surroundings to the system.
• If energy is absorbed during a reaction then the
  enthalpy of the products will be higher than that
  of the reactants.
• This means the enthalpy change will have a
  positive sign.
• This reaction will either get cooler or heat will
  need to be supplied - temperature decreases.

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Exothermic Reaction

• An exothermic reaction is one where energy is
  transferred from the system to the surroundings.
• If energy is released during a reaction then the
  enthalpy of the products will be lower than that
  of the reactants.
• This means the enthalpy change will have a
  negative sign.
• This reaction will feel hotter - temperature
  increases.

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Stability

• Reactions in chemistry tend towards products that
  are more stable.
• Stability increases as energy decreases so
  exothermic reactions increase the stability of
  the substances.
• This is why most chemical reactions that occur in
  nature are exothermic.
• Endothermic reactions usually need help in the
  form of energy to allow them to occur.

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Stability

• Remember that lower energy is more stable. We can
  compare this to standing on top of a high
  building where you have more potential energy
  than someone on the ground.
• If you fall to the ground you lose some of that
  potential energy you had on the roof but you are
  now more stable.

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Exothermic and Endothermic Reactions

• 5.1.2 State that combustion and neutralization
  are exothermic processes.

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Combustion

• Combustion is the scientific word for burning.
• Most hydrocarbons burn easily in excess oxygen.
• When they burn they produce carbon dioxide and
  water.
• Ex. CH4(g) 2O2(g) ? CO2(g) 2H2O(g)
• Combustion produces lots of heat which is
  transferred to the surroundings so it is very
  EXOthermic.

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Combustion

• When 1 mole of a substance is burned the energy
  released is called the Enthalpy Change of
  Combustion, ?Hc.
• The sign indicates this was measured under
  standard conditions in a controlled environment.
• Standard conditions for thermochemistry
  experiments are T 298K and P 101.3kPa ( 1
  atm).

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Neutralization

• Neutralization reactions involved acids and
  bases.
• If an acid and a base react completely the
  resulting solution will be pH neutral.
• The products are a salt (ionic compound) water.
• Ex. NaOH(aq) HCl(aq) ? NaCl(aq) H2O(l)
• When weak acids and bases are involved such as
  NaHCO3 then CO2(g) may be another product.
• These reactions release energy to the
  surroundings so they are EXOthermic.

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Negative ?H

• As combustion and neutralization reactions are
  always exothermic the enthalpy changes will
  always have negative values.
• You can find many ?H values in scientific
  literature ex. your textbook, the books in the
  classroom and online.

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Spontaneous Reactions

• A spontaneous reaction is one that occurs when
  the reactants are mixed without the need to be
  heated or have some other outside influence.
• Most spontaneous reactions are EXOthermic but
  there are some spontaneous endothermic reactions
  ex. Dissolving NH4Cl in water.

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Summary of Enthalpy Changes
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Enthalpy Changes

• In an exothermic reaction the products are more
  stable than the reactants so the bonds made are
  stronger than the bonds broken.
• In an endothermic reaction the products are less
  stable than the reactants so the bonds made are
  weaker than the bonds broken.

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Enthalpy Changes

• Enthalpy changes are usually written alongside
  the chemical equation for the process with a
  positive or negative sign.
• State symbols are VERY IMPORTANT as changes of
  state have their own enthalpy change values.
• Enthalpy changes are usually reported per mole so
  the units are kJ mol-1.
• If this is not the case then just kJ is used.

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Enthalpy Changes

• For example
• 2NaHCO3(s) ? Na2CO3(s) H2O(l) CO2(g)
• ?H 91.6 kJ mol-1
• The sign indicates its an endothermic reaction.

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Standard Conditions

• To compare enthalpy changes conditions must be
  the same.
• The thermochemical standard conditions are
• Temperature 25C 298K (this is room temp)
• Pressure 1 atm 101.3kPa
• Solutions have a concentration of 1 mol dm-3
• Standard conditions are sometimes indicated by
  the symbol ? or ?H? or?H
• Sometimes the temperature is included too ?H?
  298.
• The values in bold are the SI units.

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Learning Check Do NOW

• Write the equation for the formation of chlorine
  oxide, Cl2O from its elements
• What bonds are broken and what bonds are made in
  this process?
• Do the processes in 2. absorb or release energy?
• What is an enthalpy change?
• In this reaction the bonds made are less strong
  than those broken. Will the enthalpy change be
  positive or negative?
• Will this be an exothermic or endothermic
  reaction?

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Calculation of Enthalpy Changes

• 5.2.1 Calculate the heat energy change when the
  temperature of a pure substance is changed.
• Students should be able to calculate the heat
  energy change for a substance given the mass,
  specific heat capacity and temperature change
  using q mc?T.
• 5.2.2 Design suitable experimental procedures for
  measuring the heat energy changes of reactions.
  Students should consider reactions in aqueous
  solution and combustion reactions.

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Calculation of enthalpy changes

• 5.2.3 Calculate the enthalpy change for a
  reaction using experimental data on temperature
  changes, quantities of reactants and mass of
  water.
• 5.2.4 Evaluate the results of experiments to
  determine enthalpy changes. Students should be
  aware of the assumptions made and errors due to
  heat loss.

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Temperature

• Temperature is a measure of the average kinetic
  energy of the particles in a system.
• Units are K or C.
• Note there is no sign used with K!
• Heat is a measure of the total energy in a
  substance.
• TC 273 TK
• TK 273 TC

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Specific Heat Capcity

• Some substances will conduct heat and therefore
  change temperature more easily than others.
• Ex. A metal pan on a stove will become very hot
  before the water inside it does.
• A measure of how easily something changes
  temperature is called SPECIFIC HEAT CAPACITY.

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Specific Heat Capacity

• Specific heat capacity is the amount of heat
  energy required to increase 1 g of a substance by
  1K or 1C.
• It is used in the following equation
• q m x c x ?T where q is heat energy, m is mass,
  c is specific heat capacity and ?T is change in
  temperature.

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Specific Heat Capacity

• Units of specific heat capacity (c) are J g-1K-1
• Units of heat energy (q) are J or kJ
• Units of mass (m) are g or kg
• Units of temperature (T) are K or C

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Calorimetry

• Calorimetry is a method used to measure the
  enthalpy associated with a particular change.
• The temperature change of a liquid is measured
  inside a well insulated container called a
  calorimeter.
• Often a styrofoam cup is used as it has a very
  low heat capacity and is a good insulator.

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Measuring Energy Changes

• To measure the enthalpy change of a reaction that
  occurs in solution you can carry it out in a
  styrofoam (polystyrene) cup and monitor the
  temperature during the reaction.
• Styrofoam is a good insulator so the amount of
  heat lost to the surroundings will be reduced.

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Measuring Energy Changes

• Burning substances in a bomb calorimeter
  measures the temperature change in the water
  surrounding the burning item.
• This system is also well insulated to try and
  reduce heat loss.
• There may be some losses from incomplete
  combustion.

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Calorimetry

• If calorimeters made of other materials are used
  then the heat absorbed by the calorimeter must be
  added to that absorbed by the liquid
• Heat absorbed (mc?T)liquid (mc?T)calorimeter
• Calorimetry assumes no heat is transferred to or
  from the surroundings so they must be well
  insulated.
• However this is hard to achieve and is a major
  source of error in high school labs.

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Calorimetry

• Once you have calculated the energy released from
  a process then you can calculate how many kJ of
  energy were released per mole of the reactant.
  This is usually referred to as ?H or molar
  enthalpy in a test question.
• ?H q/n
• Where n represents the number of moles of the
  reactant that is reacting/burning.

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Sample Problems

• Specific heat, c of liquid water 4.18 kJ
  dm-3K-1
• 4.18J g-1K-1
• How much heat energy is required to increase the
  temperature of 10 g of nickel (c 440 J g-1K-1)
  from 50C to 70C?
• The enthalpy of combustion of ethanol (C2H5OH) is
  1370 kJ mol-1. How much heat is released when
  0.200 mol undergo complete combustion?

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Sample Problems

• 3. H2(g) 1/2 O2(g) ? H2O(l)
• ?H for the reaction above is -286 kJ mol-1. What
  mass of oxygen must be consumed to produce 1144
  kJ of energy?
• 4. Calculate the molar enthalpy change when
  excess zinc is added to 50 cm3 of a 1 mol dm-3
  solution of CuSO4. The temperature increases from
  20C to 70C when the zinc is added. Assume the
  solution has the same density as water 1.00 g
  cm-3

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Bond Enthalpies

• 5.4.1 Define the term average bond enthalpy.
• 5.4.2 Explain, in terms of average bond
  enthalpies, why some reactions are exothermic and
  others are endothermic.

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Bond Enthalpies

• All chemical reactions involve the making and
  breaking of bonds.
• The bonds in the reactants are broken which
  absorbs energy so this is an endothermic process.
• The bonds in the products form which releases
  energy so this is an exothermic process.

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Bond Enthalpy

• Bond enthalpy is defined as the energy needed to
  break one mole of bonds in gaseous molecules
  under standard conditions.
• Ex. ½ H2(g) ½ Cl2(g) ? H(g) Cl(g)
• Breaking a bond is endothermic so these values
  are always positive.
• Bond enthalpies depend on the rest of the
  molecule so values are usually averages.

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Bond Enthalpy

• The energy released when a bond is made is the
  same value as the bond enthalpy but with a
  negative sign.
• A higher bond enthalpy indicates a stronger bond.

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Bond Enthalpy

• If the bonds being broken (bonds in the
  reactants) are weaker than those being made
  (bonds in products) then the reaction will be
  exothermic.
• If the bonds being broken are stronger than those
  being made then the reaction will be endothermic.

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Bond Enthalpies

• As bond enthalpies are averages and they are only
  for gases they are not the most accurate way to
  calculate an enthalpy change but they are usually
  within about 10 of more accurate values and are
  a useful tool.
• ?Hºreaction ?BEbonds broken- ?BEbonds made
• Where BE stands for bond enthalpy

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Bond Enthalpy

• Bond enthalpy values are given in the data
  booklet. A sample is shown here. When using these
  values be careful to check if bonds are single or
  double etc.

Bond E / kJ mol-1
H-H 436
C-H 413
C-C 347
CC 612
OO 498
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Bond Enthalpy Calculations

• When water is formed from its elements, what
  bonds are broken and formed? What is the enthalpy
  change predicted by bond enthalpies?
• 2H2(g) O2(g) ? 2H2O(l)
• Bonds broken 2(H-H), 1 (OO)
• Bonds formed 4 (O-H)
• ?H Sbonds broken Sbonds formed

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Bond Enthalpy Calculations

• (2(436) 498 ) (4(464))
• (872 498) (1856)
• 1370 1856 -486 kJ
• What does this sign tell you about the reaction?
• Does this make sense?

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Practice Problem

• Find the enthalpy change when CO2(g) is formed
  from its elements using bond enthalpy values in
  the data booklet.

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Hesss Law

• 5.3.1 Determine the enthalpy change of a reaction
  that is the sum of two or three reactions with
  known enthalpy changes.
• Students should be able to use simple enthalpy
  cycles and enthalpy level diagrams and to
  manipulate equations. Students will not be
  required to state Hesss Law.

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Hesss Law

• Hesss Law is a special case of the law of
  conservation of energy.
• It states that the enthalpy change for a reaction
  will have the same value no matter how many steps
  were taken to go from reactants to products.
• Another way we can use it is to say the enthalpy
  change for the reaction is equal to the enthalpy
  of the products - enthalpy of reactants.

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Hesss Law

• Hesss Law enables us to calculate an enthalpy
  change for a reaction without carrying out the
  actual reaction.
• We do this by measuring the enthalpy change for
  other related reactions.
• Hesss Law is very useful for reactions that are
  difficult to carry out in a lab or that do not
  occur.

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Hesss Law

• For example sodium hydrogen carbonate can be
  reacted with hydrochloric acid as follows
• NaHCO3(s) HCl(aq) ? NaCl(aq) CO2(g) H2O(l)
  ?H1
• To give the same products we could carry out 2
  other reactions
• 2NaHCO3(s) ? Na2CO3(s) CO2(g) H2O(l)
  ?H2
• Na2CO3(s) 2HCl (aq) ? 2NaCl(aq) CO2(g)
  H2O(l) ?H3

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Hesss Law

• The sodium hydrogencarbonate could be heated to
  form sodium carbonate which is then reacted with
  hydrochloric acid.
• 2NaHCO3(s) ? Na2CO3(s) CO2(g) H2O(l) ?H2
• Na2CO3(s) 2HCl (aq) ? 2NaCl(aq) CO2(g)
  H2O(l) ?H3

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Hesss Law

• If these two equations are added together the
  Na2CO3 cancel out and the result is twice the
  overall equation.
• 2NaHCO3(s) 2HCl(aq) --gt 2NaCl(aq) 2CO2(g)
  2H2O(l)
• Hesss Law says that the enthalpy change for the
  two stage reaction must be equal to the single
  stage process
• 2?H1 ?H2 ?H3
• The number 2 is present in front of ?H1 as the
  equations must all be balanced.

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Hesss Law

• This can also be shown using an enthalpy cycle

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Hesss Law

• To read an energy cycle you must identify the
  reactants and products that you are being asked
  about and make sure all 3 sides are balanced.
• Then you must identify the alternative route you
  will take to get from reactants to products.

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Hesss Law

• If you are following an arrow then you add all
  those ?H values. You then subtract any ?H values
  when you go in the opposite direction of the
  arrow.

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Hesss Law

• You can also use a series of equations and
  compare them to find the unknown enthalpy change.
• For example
• S(s) 1½O2(g) ? SO3(g) ?H-395 kJ ?H1
• SO2(g) ½O2(g) ? SO3(g) ?H-98 kJ ?H2
• Calculate the standard enthalpy change, ?H for
  the reaction
• S(s) O2 ? SO2(g)

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Hesss Law

• ?H1 has the same reactants as the reaction in
  question.
• ?H2 relates SO3(g) to SO2(g) but the SO2 is on
  the opposite side of the equation from the
  reaction in question. This means we reverse the
  chemical change and the sign of ?H2.
• The equations can now be combined and the values
  added to find the ?H
• ?H -395 98 -297 kJ

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Hesss Law Activity Worksheet

• 1. Calculate ?H for the reaction
• C2H4 (g) H2 (g) ? C2H6 (g),
• from the following data.
• C2H4 (g) 3 O2 (g) ? 2 CO2 (g) 2 H2O (l) ?H
  -1411. kJ/mole
• C2H6 (g) 7/2 O2 (g) ?2 CO2 (g) 3 H2O (l) ?H
  -1560. kJ/mole
• H2 (g) 1/2 O2 (g) ?H2O (l) ?H -285.8
  kJ/mole

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Solution to 1 problem

• Reactions that were reversed or multiplied by a
  constant are shown in italics.
• 1. ?H -137. kJ
• C2H4 (g) 3 O2 (g) ? 2 CO2 (g) 2 H2O (l) ?H
  -1411. kJ
• 2 CO2 (g) 3 H2O (l) ? C2H6 (g) 7/2 O2 (g) ?H
  1560. kJ
• H2 (g) 1/2 O2 (g) ? H2O (l) ?H -285.8 kJ

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Hesss Law Worksheet

• 2. Calculate ?H for the reaction
• 4 NH3 (g) 5 O2 (g) ?4 NO (g) 6 H2O (g),
• from the following data.
• N2 (g) O2 (g) ?2 NO (g) ?H -180.5 kJ
• N2 (g) 3 H2 (g)? 2 NH3 (g) ?H -91.8 kJ
• 2 H2 (g) O2 (g) ? 2 H2O (g) ?H -483.6 kJ
• Answer
• ?H -1628. kJ